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5.1 - Exothermic and Endothermic Reactions
5.1.1 - Define the terms exothermic reaction, endothermic reaction and standard enthalpy
change of reaction
Exothermic Reaction - A reaction that causes the temperature of the surroundings to
increase.
Energy is lost, or released, in the reaction, as the enthalpy of the products is less
than the enthalpy of the reactants.
Endothermic Reaction - A reaction that causes the temperature of the surroundings to
decrease.
Energy is used up in the reaction, as the enthalpy of the products is greater than the
enthalpy of the reactants. Most endothermic reactions are not spontaneous
Standard Enthalpy Change of Reaction - The difference between the products and the
enthalpy of the reactants under standard conditions.
This is affected by temperature, pressure, concentration of solutions and physical
state. Therefore, we have to make sure that these are specified, otherwise results
are not comparable. The standard conditions of heat energy transfer are pressure
101.3kPa and temperature 298K. Only the change in H can be measured for a given
for a reaction, not H for the initial or final state of the species.
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5.1.2 - State that combustion and neutralisation are exothermic reactions
Combustion is an exothermic reaction as heat energy is released into the surroundings
Neutralisation reactions are exothermic
These are the reactions between acids and bases to form salt and water
5.1.3 - Apply the relationship between temperature change, enthalpy change and the
classification of reactions as endothermic or exothermic
When the enthalpy of the products is greater than the reactants, energy is absorbed from
the surroundings: an endothermic reaction. The temperature of the surroundings decreases
o This kind of reaction is used for ice packs in sports injuries
This is represented mathematically as the enthalpy of the products minus that of the
reactants to find the change
When the change is negative, energy is lost to the surroundings, making the reaction
exothermic
When the change is positive, energy is used from the surroundings, making the reaction
endothermic
A thermochemical equation is when the equation is written with an associated change in
enthalpy
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5.1.4 - Deduce, from an enthalpy level diagram, the relative stabilities of reactants and
products, and the sign of the enthalpy change for the reaction
Enthalpy level diagrams, or energy profile diagrams, allow us to visualise what happens to
the enthalpy of a reaction as it proceeds
The total enthalpy of the reactant species is labelled HR, the products HP
The bonds between the reactants must first be broken before they can be converted into
products. They must collide with sufficient energy for this to happen. Some of this kinetic
energy is converted into vibrational energy, which overcomes the bonds when it reaches a
certain magnitude. Only then can the reaction be initiated
This amount of energy is known as the activation energy, Ea, seen as a hump on the
diagram
The distance between HR and the top of the hump is the total amount of activation energy
required. This is always a positive value
When bonds are made to form the products, this always releases energy, causing the
diagram to drop at HP
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We can also deduce the relative stabilities of the reactants and products
The higher the enthalpy, the less stable the substance
In exothermic reactions, the products have a lower enthalpy than the reactants, and so are
more stable, and vice versa for endothermic reactions. Large activation energies indicate
strong bonds in the reactants, i.e.
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Other reactions have small activation energies. The reaction with white phosphorus below is
exothermic. White phosphorus is highly reactive and dangerous, so the less reactive red
phosphorus is used in safety matches
The instability of the reactants makes it easier for their bonds to be broken, which is why
the activation energy is so low.
Here are a few more reactions you need to be aware of:
Ionisation – Endothermic
Change of State: Melting, Evaporation and Sublimation - Endothermic
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Change of State: Vaporisation and Condensation – Exothermic
Remember that bond breaking is endothermic and bond making is exothermic
Summary
Type of Thermochemical
Reaction
Exothermic Endothermic
Enthalpy change Products < Reactants Products > Reactants
ΔH negative positive
Change in surrounding
temperature
increases decreases
Enthalpy level diagram
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5.2 – Calculation of Enthalpy Changes
5.2.1 - Calculate the heat energy change when the temperature of a pure substance is
changed
Specific Heat Capacity – Symbol c – The amount of heat energy required to raise the
temperature of 1.00g of the substance by 1.00°C or 1.00K
The specific heat capacity for some common substances is shown below
Substance Specific Heat
Capacity (J °C-1 g-1)
Water 4.18
Aluminium 0.889
Ammonia 2.06
Copper 0.386
Ethanol 2.46
Gold 0.128
Iron 0.448
The equation to calculate the heat energy change is
q = heat energy change – J
m = mass of the substance – g
c = specific heat capacity of the substance – J °C-1 g-1 or J K-1 g-1
ΔT = change in temperature - °C or K
For the equation to work, all these figures must relate to the one substance
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5.2.2 - Design suitable experimental procedures for measuring the heat energy changes of
reactions
Here, you need to know two types of experiments. One involves a reaction in aqueous
solution and the other involves combustion reactions.
A simple Calorimetry is the most appropriate method here. In this type of procedure, the
heat energy released is measured
The heat transfers to the can of water holding the reactants; however this is an incomplete
transfer. The calculations of enthalpy change are based on the temperature change of the
water and the mass of the fuel.
Reactions in Aqueous Solutions
These reactions can be done in a simple calorimeter, such as a Styrofoam cup. The volume
of water must be accurately known for accurate calculations.
There will always be error in this measurement as heat is lost to the surroundings, as well as
the can
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Combustion Reactions
For these reactions, we measure the amount of heat produced by the combustion of one
mole of the fuel. Note that this is not necessarily the same as the enthalpy of combustion.
The heat of combustion for the reaction below is actually 5464 kJ mol-1, however the
enthalpy is twice this because there are two moles of octane involved.
When the fuels are used in the form of complex mixtures, the heat of combustion is
expressed as kJ g-1 or MJ dm-3
Some examples of the heat of combustion for certain compounds are given below
Hydrocarbon Heat of Combustion
kJ mol-1 kJ g-1
Methane 889 55.6
Ethane 1557 51.9
Propane 2217 50.4
Butane 2874 49.6
Octane 5464 47.9
Common Fuels Heat of Combustion
kJ g-1
MJ dm-3
Liquid Hydrogen 142 286
Ethanol 29.7 19.6
Petrol 47.3 29.0
Diesel 44.8 32.2
LPG 51 22.2
Wood 15 n/a
Brown, dried coal 25 n/a
Black, dried coal 32 n/a
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5.2.3 - Calculate the enthalpy change for a reaction using experimental data on
temperature changes, quantities of reactants and mass of water
In thermochemical equations, the value for ΔH is always given for the equation exactly as it
is written. Therefore, if you double the equation, then ΔH is also doubled. Also, if the
reaction is reversed [swapping the products and the reactants], then the sign changes.
Compare the equations below
For example:
Calculate the amount of energy released when 500cm3 of methane gas at STP reacts with
excess air according to the equation
Therefore, the energy released in this combustion reaction is 19.9 kJ
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5.2.4 - Evaluate the results of experiments to determine enthalpy changes
Basically what you need to do here is take the results from an experiment, put them into the
equation, and do the appropriate calculations.
Once you have done this, you also need to be able to consider the problems or errors with
the results. These should include:
In combustion reactions, heat is lost to the surroundings
There is a great deal of error in this method
o Does not allow for the heat energy absorbed by the can or
polystyrene cup
Therefore, the calculated heat energy change is inaccurate
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5.3 - Hess’s Law
5.3.1 - Determine the enthalpy change of a reaction that is the sum of two or three
reaction with known enthalpy changes
The energy difference between two states is
independent of the route between them
i.e. The heat evolved or absorbed in a
chemical process is the same, whether
the process takes place in one or
several steps
Energy can be neither created nor destroyed, it can only change state. This law can be used
to determine the enthalpy of a reaction by manipulating known equations that could be
used as a reaction pathway to the desired reaction
i.e. To determine the enthalpy for , we can use
and
The reactions are added. They are all put into an enthalpy cycle to show the relationship
between them.
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5.4 - Bond Enthalpies
5.4.1 - Define the term average bond enthalpy
Average Bond Enthalpy - The amount of energy required to break one mole of bonds in the
gaseous state averaged across a range of compounds containing that bond.
For example, in determining the bond enthalpy of C-H:
5.4.2 - Explain, in terms of average bond enthalpies, why some reactions are exothermic
and others are endothermic
Bond breaking
This is an endothermic process, as energy is absorbed to break the bonds between the
atoms.
Bond making
This is an exothermic process, as energy is released when the bonds form between the
atoms.
If more energy is released in the formation of bonds in the products that was required to
break the bonds of the reactants, then the reaction is exothermic. If more energy is
required to break the bonds of the reactants than is released in the formation of bonds in
the products, then the reaction is endothermic